Download Periodic Properties

Periodic Properties
3.2 Physical Properties
3.2.1
Define the terms first ionization energy and electronegativity.
3.2.2
Describe and explain the trends in atomic radii, ionic radii, first ionization
energies, electronegativities and melting points for the alkali metals (Li  Cs) and the
halogens (F  I).
3.2.3
Describe and explain the trends in atomic radii, ionic radii, first ionization
energies and electronegativities for elements across period 3.
3.2.4
Compare the relative electronegativity values of two or more elements based
on their positions in the periodic table.
.1
Periodic Properties
• Elements show gradual changes in certain physical
properties as one moves across a period or down a
group in the periodic table. These properties repeat
after certain intervals. In other words they are
PERIODIC
Periodic properties
include:
-- Ionization Energy
-- Electronegativity
-- Electron Affinity
-- Atomic Radius
-- Ionic Radius
.2
Periodic Properties
Periodic properties include:
-- Ionization Energy
-- Electronegativity
-- Electron Affinity
-- Atomic Radius
-- Ionic Radius
• All these properties and trends are a result of:
Effective Nuclear Charge (Eff. N. Charge)
This refers to how effectively the protons do (or do not)
pull on their own electrons and the electrons of
neighboring atoms.
.3
Periodic Properties
• As you move across the periodic table (in a period),
you increase both Protons (+) and electrons (-).
*but think of their locations!!!
Protons are centrally located (in one point).
Electrons are spaced out surrounding the
nucleus.
.4
• Metals lose electrons more easily than
nonmetals.
.5
• Nonmetals lose electrons with difficulty. (They
like to GAIN electrons).
.6
Trends in Ionization Energy
Ionization energy is the energy required to
remove an electron from an atom
• Metals lose electrons more
easily.
• Nonmetals gain electrons.
• Ionization energy increases
across a period because the
pull increases. (called what?)
.7
Trends in Ionization Energy
• The ionization energy (IE) is
highest at the top of a group. IE
decreases as the atom size
increases.
• This results from an effect
known as the Shielding Effect
.8
The Electron Shielding Effect
• Electrons between
the nucleus and
the valence
electrons repel
each other
making the atom
larger.
.9
Ionization Energies of the
Representative Groups
.10
Ionization Energies are Periodic
.11
Electronegativity
Electronegativity is a
measure of the
ability of an atom
attract electrons IN
A CHEMICAL BOND.
This concept was first proposed by Linus Pauling
(1901-1994). He later won the Nobel Prize for his
efforts.
.12
Periodic Trends: Electronegativity
• In a group: Atoms with fewer
energy levels can attract electrons
better (less shielding). So,
electronegativity increases UP a
group of elements.
• In a period: More protons, while
the energy levels are the same,
means atoms can better attract
electrons. So, electronegativity
increases RIGHT in a period of
elements.
.13
Trends in Electronegativity
Electronegativity increases across
a period and up a group
.14
Electronegativity
.15
Electronegativity
.16
Electron Affinities
The Electron affinity of an atom is the energy change when
an electron is added to the neutral atom to form a negative
ion.
*This property can only be measured in an atom in gaseous
state.
X + e− → X−
.17
Electron Affinities
.18
Electron Affinities Are Periodic
• Electron Affinity vs. Atomic Number
.19
The Electron Shielding Effect
• Electrons between
the nucleus and
the valence
electrons repel
each other
making the atom
larger.
.20
Atomic
Radius
• The radius increases on going down a group.
• Because electrons are added further from the
nucleus, there is less attraction. This is due to
additional energy levels and the shielding effect.
Each additional energy level “shields” the electrons
from being pulled in toward the nucleus.
• The radius decreases on going across a period.
.21
Atomic Radius
• The radius decreases across a period owing to
increase in the positive charge from the protons.
• Each added electron feels a greater and greater +
charge because the protons are pulling in the same
direction, whereas the electrons are scattered.
Large
All values are in
nanometers
Small
.22
Atomic Radius
.23
Atomic Radius
.24
Trends in Ion Sizes
Radius in pm
.25
Cations
Cations (positive ions) are smaller than
their corresponding atoms
.26
Ionic Radius
+
Li
0.152 nm
3e and 3p
Li + ,
0.078 nm
2e and 3 p
Forming a
cation.
• CATIONS are SMALLER than the atoms from which
they come.
• The electron/proton attraction has gone UP and so
the radius DECREASES.
.27
Ionic Radius for Cations
Positive ions or
cations are smaller
than the
corresponding
atoms.
Cations like atoms
increase as one
moves from top to
bottom in a group.
.28
Anions
Anions (negative ions) are larger than
their corresponding atoms
.29
Ionic Radius-Anions
F 0.064 nm
9e- and 9p+
F- 0.133 nm
10 e- and 9 p+
Forming an anion.
• ANIONS are LARGER than the atoms from
which they come.
• The electron/proton attraction has gone
DOWN and so size INCREASES.
• Trends in ion sizes are the same as atom sizes.
.30
Ion Sizes
Does the size go up or
down when gaining an
electron to form an anion?
.31
Ionic Radii for Anions
Negative ions or
anions are larger
than the
corresponding
atoms.
Anions like atoms
increase as one
moves from top to
bottom in a group.
.32
Ionic Radius for an
Isoelectronic Group
Isoelectronic ions
have the same
number of electrons.
The more negative an
ion is the larger it is
and vice versa.
.33
Summary of Periodic Trends